Chem Exam 3

avg atomic mass

avg atomic mass = the sum of each isotope mass times its abundance

diatomic elements

H, N, Fl, O, Cl, I, Br

finding empirical formula and molecular formula from % composition

• turn %s into g (assume sample is 100g)

• find mols of each element

• divide by smallest number of mols

• round to nearest integer

• divide mass of molecular compound by mass of empirical formula to find factor of each element

equations where one reactant is in excess

• find how many mols of each reactant

• find possible mass of product from limiting reactant

• find how much of the other reactant is needed for that product

• subtract that mass from total mass of excess reactant to find unreacted mass

% yield

% yield = actual yield / theoretical yield

determining MM of unknown reactant

• calculate mols of product

• calculate proportionate mols of reactant

• divide mass of reactant by mols

Solubility Rules

• Soluble - alkali metals, ammonium, (NO3)-, (CH3CO2)-, Cl/Br/I

  • not with Cu+, Ag+, Pb 2+

• Insoluble - (OH)-, (S) 2-, (CO3)2-, (PO4)3-, (SO4)2- with a main group cation

  • not with, alkali metal, NH4, alkali earth metal, Mg2+ or 2+ transition metal

Strong Acids - break apart and ionize, high conductivity, covalent

HCl, HBr, HI, HNO3, HClO4, H2SO4

Strong Base - break apart and ionize, high conductivity, ionic

LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Electrolytes - conduct electricity

Strong - soluble compounds, strong acids/bases, ions

Weak - weak acids/bases

non-electrolyte - insoluble compound

Oxidation Numbers

• element in naturally occurring form: 0

• charged monatomic ions: their charge

• commands: sum of ox# = 0

  • oxygen: -2

  • hydrogen: +1

  • halogens: -1

OIL RIG

if ox# increase, electrons are lost, species is oxidized, reducing agent

if ox# decrease, electrons are gained, species is reduced, oxidizing agent

types of reactions

• precipitation (forms a solid)

• acid-base (reactants are an acid and a base)

  • neutralization

  • gas formation

• oxidation reduction (elements change ox#)

  • combustion

  • single displacement

dilution equation

M1V1 = M2V2

conversions

1kcal = 4.184 kJ

0 C = 273.15 K

energy and enthalpy equations

• KE = 0.5mv² (kg, m/s)

• Δ𝐸=𝑤+𝑞 (kJ)

  • internal energy = work + heat

  • work done by system and heat released will be negative

• w=-pΔV (atm, L)

  • work = -(pressure) (change in volume)

• 𝑞=𝑚𝑐Δ𝑇 (J or cal, g, C or K) *c and T units must match

  • heat absorbed/released = (mass)(specific heat capacity)(change in temp)

  • specific heat of water = 4.184
    

path vs state functions

• path - depends on how system got from one state to another

  • heat, work

• state - depends only on current state of system

  • T, P, V, ΔE, ΔH

calculate ΔH_rxn by manipulating equations

• multiply and divide, or flip equations so that products and reactants cancel

• add enthalpies

calculate ΔH_rxn from H_f

• find the enthalpy of formation of each species from chart

• products - reactants

IMFs

• Intermolecular

  • London Dispersion

  • Dipole-Dipole

  • H-Bonding (hydrogen must be bonded to NOF, other NOF must have lone pair)

• Intramolecular

  • covalent

  • metallic

  • ionic

ideal gas law

PV=nRT

(atm, L, mol, 0.0821, K)

combined gas law

(P1V1/n1T1) = (P2V2/n2T2)

density and mass equations

• MM = mRT/PV

• MM= dRT/P

partial pressures

the total pressure of a gaseous mixture is equal to the sum of the partial pressures

diffusion vs effusion

diffusion - gases move through space until they’ve mixed completely

effusion - gases move through a small pinhole

properties of substances with strong IMFs

• high boiling point

• low vapor pressure

• high surface tension

• high viscosity

phase diagram

• triple point - all three phases can exist at once

• critical point - distinction between gas and liquid is lost

spontaneous reaction

displacement reaction will occur if lone cation is move reactive than cation in compound